Question

Why is $d_{x^2-y^2}$ in $dsp^2$ hybridization and not the other d orbitals?

• A. The $d_{x^2-y^2}$ orbital is used because its lobes are directed along the x and y axes, which align with the positions of ligands in a square planar geometry, allowing for strong sigma bond formation.
• B. The $d_{z^2}$ orbital is used because it is oriented along the z-axis, which is suitable for square planar geometry.
• C. The $d_{xy}$ orbital is used because its lobes are oriented between the x and y axes, providing better overlap in square planar arrangements.
• D. Any d orbital can be used in $dsp^2$ hybridization as their orientations are interchangeable.

Answer 1

Answer: (A)

The $d_{x^2-y^2}$ orbital is used because its lobes are directed along the x and y axes, which align with the positions of ligands in a square planar geometry, allowing for strong sigma bond formation.

Answer 2

In $dsp^2$ hybridization, which leads to a square planar geometry, four hybrid orbitals are formed and directed towards the corners of a square in a plane (e.g., the xy plane). The $d_{x^2-y^2}$ atomic orbital has its lobes oriented precisely along the x and y axes. This specific orientation allows for maximum head-on overlap with ligand orbitals positioned along these axes, facilitating the formation of strong sigma bonds. The other d orbitals, such as $d_{xy}$ (lobes between axes), $d_{yz}$ and $d_{xz}$ (lobes in perpendicular planes), and $d_{z^2}$ (lobes along the z-axis), are not optimally aligned for direct overlap with ligands in this square planar arrangement.